# Covalent Bonds

Generally, covalent bonds between atoms with identical (or almost identical) electronegativity will be nonpolar covalent bonds. Covalent bonds between atoms with different electronegativity will be polar covalent bonds. If the difference is big enough, the compound will have an ionic bond. Molecules with polar covalent bonds that do not cancel will have a net dipole moment.

A covalent bond is a chemical bond that forms when valence electrons are shared between atoms. Nonpolar covalent bonds form between two identical elements or between elements with identical affinities for electrons. A nonpolar covalent bond is a covalent bond in which the nuclei of the bonded atoms exert equal or nearly equal force on the shared electrons. For example, H2 contains nonpolar covalent bonds because the hydrogen (H) atoms are identical and they attract their shared electrons with identical strength. The tendency of an atom to attract electrons toward itself when forming bonds is its electronegativity.

Atoms with different electronegativities form polar covalent bonds. A polar covalent bond is a covalent bond in which the electron density is more localized on one end of the bond. One end is slightly positive, and one end is slightly negative. In these molecules, the shared electrons in the covalent bond pull toward the atom with the higher electronegativity. For example, hydrogen fluoride has a polar covalent bond because fluorine is more electronegative than hydrogen, so it pulls the electrons toward itself, and the distribution of electrons in the bond is polarized. A covalent bond becomes more polar as the difference between the electronegativities of its atoms increases.
Generally, electronegativity increases from left to right across the periodic table and decreases down the columns. Fluorine is the most electronegative element, followed by oxygen.

#### Electronegativity Trends

The polarity of a covalent bond can be quantified by its dipole moment. A dipole moment is a vector quantity that defines the extent of the charge on either side of a polar covalent bond, with the direction that points from the positive side of the bond toward the negative side. A dipole forms whenever opposite charges are separated from each other, and a dipole moment is calculated by multiplying the amount of the charge by the distance between the centers of the charge. Dipole moments are reported in debye (D), and a larger dipole moment indicates a more polar bond. Dipole moments also have direction, with the dipole going toward the element with greater electronegativity.

### Selected Bond Dipole Moments of Hydrogen

Bond Dipole Moment
(debye)
${\rm{H{-}F}}$ 1.7
${\rm{H{-}Cl}}$ 1.1
${\rm{H{-}Br}}$ 0.8
${\rm{H{-}O}}$ 1.5

The dipole moment of a covalent bond with hydrogen is related to the difference in electronegativity of the hydrogen atom and the other atom in the covalent bond.

### Selected Bond Dipole Moments of Carbon

Bond Dipole Moment
(debye)
${\rm{C{-}F}}$ 1.4
${\rm{C{-}O}}$ 0.7
${\rm{C{-}N}}$ 0.4

The dipole moment of a covalent bond with carbon is related to the difference in electronegativity of the carbon atom and the other atom in the covalent bond.

Molecules containing multiple polar covalent bonds may have a molecular dipole moment, or the polar bonds may cancel each other and result in a molecule with no dipole moment. For example, carbon dioxide contains two polar double bonds between the carbon and oxygen atoms, but the geometry (linear) of the molecule causes the dipoles to cancel each other out. By contrast, methylene chloride (tetrahedral) has a dipole moment of 1.62 D because of the different polarities of the hydrogen-carbon bond compared with the carbon-chlorine bonds.